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The Importance of Chemical Reactions in the Charging Process of Lithium-Sulfur Batteries

Berger, A.; Freiberg, A.T.S.; Siebel, A.; Thomas, R.; Patel, M.U.M.; Tromp, M.; Gasteiger,H.A.; Gorlin, Y.Published in:Journal of the Electrochemical Society

DOI:10.1149/2.0181807jes

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Citation for published version (APA):Berger, A., Freiberg, A. T. S., Siebel, A., Thomas, R., Patel, M. U. M., Tromp, M., ... Gorlin, Y. (2018). TheImportance of Chemical Reactions in the Charging Process of Lithium-Sulfur Batteries. Journal of theElectrochemical Society, 165(7), A1288-A1296. https://doi.org/10.1149/2.0181807jes

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Download date: 16 Apr 2020

A1288 Journal of The Electrochemical Society, 165 (7) A1288-A1296 (2018)

The Importance of Chemical Reactions in the Charging Process ofLithium-Sulfur BatteriesAnne Berger,1,∗ Anna T. S. Freiberg,1,∗ Armin Siebel, 1,∗ Rowena Thomas,1Manu U. M. Patel,1 Moniek Tromp,2 Hubert A. Gasteiger,1,∗∗ and Yelena Gorlin 1,∗∗∗,a,z

1Chair of Technical Electrochemistry, Department of Chemistry and Catalysis Research Center, Technische UniversitatMunchen, Garching, Germany2Sustainable Materials Characterisation, Van’t Hoff Institute for Molecular Sciences, University of Amsterdam,Amsterdam, Netherlands

The underlying mechanism of lithium-sulfur batteries is still not fully established because it involves a series of both chemical andelectrochemical reactions as well as the formation of soluble polysulfide intermediates. To improve the mechanistic understanding oflithium-sulfur batteries, this study investigates chemical reactions between the Li2S cathode and more oxidized sulfur species, suchas S8 and polysulfides, during the electrochemical charge of the battery. By combining the electrochemistry with X-ray absorptionspectroscopy, we show that chemical reactions and, in particular, the resulting accumulation of solution species in the electrolyteare essential to oxidize Li2S at a low overpotential. Additionally, by efficiently separating the anode and cathode compartments ofa battery with a lithium ion-exchanged Nafion interlayer, we establish the adverse effect of the anode on the buildup of solutionintermediates. In the absence of the interlayer, polysulfide intermediates can diffuse through the separator and react at the anode’ssurface, while the addition of the interlayer allows the intermediates to accumulate in the separator of the cathode compartment andfacilitate the oxidation of Li2S.© The Author(s) 2018. Published by ECS. This is an open access article distributed under the terms of the Creative CommonsAttribution 4.0 License (CC BY, http://creativecommons.org/licenses/by/4.0/), which permits unrestricted reuse of the work in anymedium, provided the original work is properly cited. [DOI: 10.1149/2.0181807jes]

Manuscript submitted February 1, 2018; revised manuscript received April 5, 2018. Published May 1, 2018.

The transportation sector is going through a drastic change, asbattery powered vehicles gain increasing acceptance. However, thepresently used lithium ion (Li-ion) battery technology may not besufficient in terms of energy density, cost efficiency, and safety to ful-fill the necessary requirements for a widespread utilization of batteryelectric vehicles (BEVs).1 When both the technical and the economicconsiderations are combined, current projections suggest that BEVsusing state-of-the-art Li-ion batteries may not be able to exceed arange of 200 miles for the mid-size car market.2 Therefore, researchon post Li-ion technologies such as lithium-sulfur (Li-S) batteries havereceived increasing attention. As a cathode material, sulfur offers atheoretical gravimetric capacity of 1675 mAh/g, which is 6 times thetheoretical gravimetric capacity of NMC (Lithium Nickel Cobalt Man-ganese Oxide, 278 mAh/g), a widely used Li-ion intercalation materialfor BEVs.3 Additionally, sulfur is a promising material as it is abun-dant, inexpensive and non-toxic.4 To date, however, Li-S batteries arefar from being commercialized except for niche applications,5 as Li-Sprototypes can barely compete with state-of-the-art Li-ion technolo-gies. Their high theoretical energy density is decreased by the need ofa conductive carbon support to compensate for the insulating nature ofsulfur and the volume change during cycling.4 Additionally, the real-ization of a lithium metal anode is problematic due to safety concerns.A recent evaluation estimates that lithium-sulfur batteries might, re-alistically, not exceed the energy density of lithium-ion batteries butbe advantageous in terms of lower cost and reduced environmentalimpact.2,6

Nevertheless, as the commercial realization of Li-S batteries wouldaccelerate the progress in lightweight energy storage, it is importantto improve the cell chemistry. Instead of the intercalation mechanismin Li-ion batteries, Li-S batteries undergo a conversion reaction toform Li2S during discharge and S8 during charge through a series ofboth chemical and electrochemical reactions. During the operationof the battery, intermediate polysulfide species dissolve in the elec-trolyte and diffuse throughout the battery. These soluble intermediateshave been held responsible for major drawbacks of the system, most

∗Electrochemical Society Student Member.∗∗Electrochemical Society Fellow.

∗∗∗Electrochemical Society Member.aPresent address: Research and Technology Center, Robert Bosch LLC, California,USA.

zE-mail: yelena.gorlin@tum.de

prominently the shuttle mechanism.7–9 Therefore, many groups arefocusing on developing cathode materials that can encapsulate poly-sulfide species.10–15 Extensive research effort is also emerging in thearea of mechanistic understanding, as the determination of the exactsequence of steps occurring during the discharge and charge of thebattery would help in the design of improved electrode structures andelectrolyte solutions.

To increase the mechanistic understanding of Li-S batteries, vari-ous research groups have used both electrochemical and spectroscopictechniques, such as X-ray absorption spectroscopy (XAS),16–19 X-raydiffraction (XRD),17,20–23 and UV/Vis.24,25 In general, the studies agreethat during the discharge, the voltage curve shows two plateaus at ap-proximately 2.4 V and at 2.1 V vs. lithium metal, which are connectedby a transition zone. Shen et al. link the upper plateau to the reductionof sulfur to S8

2− and the transition zone to the reduction of S82− to

lower order polysulfides. The lower plateau is attributed to the forma-tion of Li2S.26 The idea of successive reduction steps is supported byCanas et al., who were able to distinguish via XRD the phase changefrom crystalline S8 to crystalline Li2S via a non-crystalline phase,which they attribute to a formation of soluble polysulfides.21 In addi-tion to this stepwise approach, many groups are also including multipledisproportionation reactions in their proposed mechanisms.27–29 Oneof the first to stress this necessity were Lu et al. In their study, theauthors employed a two-compartment cell, in which the cathode andanode compartments were separated by a glass membrane and a smallconcentration of S8 was dissolved in the electrolyte of the cathodecompartment. In this setup, the theoretical number of 16 electrons perS8 could be extracted at low but not high current rates, which was ratio-nalized through the inclusion of chain-growth and disproportionationsteps in the S8 reduction mechanism.27

Although most studies assume that the charging pathway is sim-ilar to the reverse of the discharging steps,21 certain differencesare recognized. Most prominently, in contrast to the two plateausobserved during the discharge, only one charging plateau is oftenseen in the most common electrolyte based on a 1:1 vol:vol mix-ture of 1,3-dioxolane (DOL) and 1,2-dimethoxyethane (DME).30

Marinescu et al. attributed the observation of one single chargingplateau to the limiting dissolution of Li2S.31 A similar mechanismis supported by Gorlin et al. who have proposed that Li2S reactschemically to form polysulfide species, which are subsequently elec-trochemically oxidized directly to S8.32 Alternatively, Yang et al. andCuisinier et al. have proposed that the charging mechanism involves

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Journal of The Electrochemical Society, 165 (7) A1288-A1296 (2018) A1289

several electrochemical reactions and therefore, several charging volt-age plateaus.19,33

Most proposed discharge and charge mechanisms include bothelectrochemical and chemical steps, but only few stress the impor-tance of the chemical reactions to the operation of Li-S batteries. Webelieve that the chemical reactions are especially important duringthe charge, a process that requires oxidation of insulating and insol-uble Li2S. In the following study we create conditions, which facili-tate chemical reactions between Li2S and solution species, by eitheradding an S8 containing layer or a polysulfide containing catholyte to aLi-S battery assembled in a discharged state. Furthermore, we applya lithium ion-exchanged Nafion interlayer to isolate the chemical re-actions occurring at the cathode from those occurring at the anode.34

These experimental conditions reveal the crucial influence of solutionspecies and the resulting chemical processes to the electrochemicalbehavior of a Li-S battery.

Although this work focuses on the charging process, the com-monly used terminology for the discharging sequence is used, whichdesignates the lithium electrode as the anode and the sulfur-containingelectrode as the cathode.

Experimental

Electrode preparation.—Li2S-electrodes.—The electrodes wereprepared inside an argon-filled glove box (MBraun; < 1 ppm H2Oand < 1 ppm O2) by weighing Li2S powder (99.98% trace metal ba-sis, Sigma-Aldrich), Vulcan carbon (XC-72, Tanaka Kikinzoku Ko-gyo) and polyvinylidene difluoride (PVDF, HSV900, Kynar) in aweight ratio of 6:3:1. The PVDF powder was dissolved in N-Methyl-2-pyrrolidone (NMP, anhydrous, 99.5%, Sigma-Aldrich) giving a so-lution of 25 mgPVDF per mLNMP. The slurry was mixed in a sealedcontainer using a planetary centrifugal vacuum mixer (Thinky). Atfirst, only carbon and Li2S were mixed together, and later PVDF dis-solved in NMP was added in successive steps. The resulting slurrywas coated on 18 μm thick aluminum foil using a 250 μm gap bar.The coating was dried overnight inside the glove box. The resultingloading was 1.5 mgLi2S/cm2. The electrodes were punched using asquare punch of the dimensions of 10 mm. Afterwards, they weredried under dynamic vacuum in a glass oven (Buchi, Switzerland) at110◦C for 10 hours. The entire procedure was performed inside theglove box or in tightly sealed vessels, preventing all contact with air.

S8-interlayer.—An S8/Vulcan carbon composite was synthesizedfollowing a previously published procedure.18,35 First, 1.5 g Vulcancarbon (XC-72, Tanaka Kikinzoku Kogyo) was dispersed in 120 mLof a 1 M sodium thiosulfate solution (Na2S2O3 · 5H2O, 99.5%, Sigma-Aldrich) for 20 minutes using a high-power sonifier. Then, 250 mL of1 M nitric acid (HNO3, ACS reagent, Sigma-Aldrich) was added tothis solution, and the mixture was sonicated for 20 minutes. In the finalstep, the suspension was washed several times, and the composite wasobtained via filtration. Prior to its use in the coating, the composite wasdried overnight under ambient conditions and further dried in a glassoven (Buchi, Switzerland) at 80◦C for 72 hours. The sulfur contentwas measured via thermogravimetric analysis (Mettler Toledo) to be68% by weight.

The sulfur interlayers were prepared by mixing the compositewith PVDF at a 9:1 weight ratio. In a sequential way, NMP was addedto give a ratio of 7 mL NMP/gPowder. The slurry was coated onto apolyolefin separator (Celgard H2013) and the electrodes were punchedin the required dimension of a square of 10 mm. The electrodes weredried over silica gel at 60◦C and static vacuum in a sealed glass oven.The resulting loading was 2.5 mgS/cm2.

Ion exchange of Nafion membrane.—Nafion HP membrane (IonPower) was ion exchanged by sequentially boiling membrane sheetsin several solutions. To activate the acid groups, the sheets were boiledin 1 M HNO3 (ACS reagent, Sigma-Aldrich) for 3 hours. Afterwards,the membrane sheets were washed and boiled for another 30 minutes indeionized water. Then, a saturated lithium hydroxide (LiOH) solution

was prepared (130 gLiOH/L) and the membranes were set to boil in apolypropylene beaker in this solution for another 6 hours. As a finalstep, the membranes were washed in warm deionized water, driedovernight, and cut into appropriate dimensions (12 mm × 14 mm).Prior to use in the battery, the membranes were dried under dynamicvacuum in a glass oven (Buchi, Switzerland) at 130◦C for an additional10 hours. The resulting lithiated Nafion membranes will be referredto as Li+-Nafion.

Cell assembly and cycling.—For all experiments, our custom-builtcell design,18 which is compatible with X-ray absorption spectroscopy(XAS) measurements, was used. A lithium metal foil (99.9% purity,450 μm, Rockwood Lithium, USA) served as the anode. Anode andcathode were separated by one layer of a 260 μm thick glass-fiberseparator (GF, glass microfiber filter 691, VWR) of the dimensions of11 mm x 13 mm. The separator was soaked with 80 μL of electrolyte.Unless stated otherwise, the electrolyte was a mixture of 1,3-dioxolane(DOL, anhydrous, 99.8%, Sigma-Aldrich) and 1,2-dimethoxyethane(DME, anhydrous, 99.8%, Sigma-Aldrich) (1:1 vol:vol) with the ad-dition of 1 M lithium perchlorate (LiClO4, battery grade, 99.99% tracemetal basis, Sigma-Aldrich) and 0.5 M lithium nitrate (LiNO3, 99.99%trace metal basis, Sigma-Aldrich). The same amount of LiNO3 wasused in all of the experiments to facilitate proper formation of a solid-electrolyte-interphase layer on the lithium metal.36 The solvents weredried for at least 72 hours over Sylobead MS 564C zeolites (3 Å,Grace Division). LiClO4 and LiNO3 powders were dried at 150◦Cunder dynamic vacuum for 72 hours. For the experiments with the S8

interlayer, an additional 40 μL of electrolyte was added. In the ex-periments that used Li+-Nafion, one separator was applied on eitherside of the membrane and 80 μL of electrolyte was added to eachseparator.

Li2S8 and Li2S4 solutions were prepared by adding Li2S and S8

powders in a stoichiometric amount to give the nominal polysulfidecomposition in DOL:DME (1:1 vol:vol). The desired polysulfide con-centration was achieved by varying the volume of the solution. S8

powder (Sigma Aldrich, 99.98%) was dried under ambient conditionsat 75◦C overnight and Li2S powder was used as received (Sigma-Aldrich, 99.98% trace metal basis). The solutions were stirred andheated to 60◦C to achieve a complete dissolution.

The assembled cells were connected to a potentiostat (Bio-LogicSAS, France) and cycled in a climatic chamber at 25◦C. The timebetween cell assembly and the start of the electrochemical experimentswas 0.5-1 hour, unless stated otherwise. The current for the galvano-static charge of the cell was set to a C-rate, ranging from C/2 to C/10,based on a theoretical gravimetric capacity of 1165 mAh/gLi2S. Thescope of this paper only covers the 1st charging sequence at constantcurrent of an electrode with Li2S as starting material. We note thatthe primary focus of our work was on isolating the chemical reactionsoccurring at the Li2S electrode from those occurring at the anodethrough the introduction of ion-exchanged Nafion interlayer. For thispurpose, all our experiments utilized the same electrolyte (added inexcess); our experiments, however, did not investigate the role of thetype of electrolyte solvent or salt, the role of used salt concentration,or the role of used electrolyte to active material ratio.

X-ray absorption spectroscopy measurements.—Sulfur K-edgeXAS measurements were performed at the PHOENIX beamline ofthe Swiss Light Source (SLS, Paul Scherrer Institut, Villigen, Switzer-land) and at beamline 14-3 of the Stanford Synchrotron RadiationLightsource (SSRL). At the SLS, the optimal setup for operando mea-surements with Li+-Nafion was determined, while the actual spectraincluded in this paper were taken at the SSRL. Beamline 14-3 is abending magnet station capable of an energy range between 2100and 5000 eV. The samples were placed inside a holder on a Newportsample station with a submicron positioning accuracy. Measurementswere carried out under a helium atmosphere and the spectra wererecorded with a Vortex silicon drift detector. The spot size was fo-cused to 15 μm in one direction (perpendicular to the cell stack) and

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A1290 Journal of The Electrochemical Society, 165 (7) A1288-A1296 (2018)

defocused to 200 μm in the other direction. An 8 μm Kapton foil(Multek) aluminized with a 100 nm layer served as an X-ray win-dow. X-ray absorption fluorescence mapping was used to visualizeand verify the operando cell alignment. The details of the measure-ment have been described previously.37 In short, fluorescence mea-surements were performed by fixing the energy of the incident X-raysto 2500 eV and rastering the irradiated spot across a defined area bymoving the sample stage in 20 μm increments. The location of sulfurwas detected by setting the region of interest (ROI) on the detectorto 2220–2470 eV. After verifying cell alignment, operando X-ray ab-sorption spectroscopy measurements were performed. The step sizefor X-ray absorption spectra was 0.2 eV in the region of interest be-tween 2460–2500 eV. To allow for better normalization an additional36 eV were scanned with a step size of 2 eV. This protocol led to a dataacquisition time of 9 minutes for one spectrum. For operando exper-iments, the spectra were recorded continuously, alternating betweenthe electrode and the separator locations. The collected data were di-vided by the intensity of the incoming X-rays to obtain raw spectra, inwhich the edge step provides a measure of the relative concentrationof sulfur species. When necessary, the spectra were also normalizedto an edge step of 1 using the Athena software package. The energyscale was calibrated to the peak position of an S8 reference spectrumat 2472 eV.38 As in our previous work, we present the raw spectrawhen the focus is on monitoring the relative concentration of sulfurspecies, and normalized spectra when the focus is on identifying thetype of sulfur species.32

In addition to fluorescence mapping, the spatial resolution wasmade certain by measuring an XAS spectrum at the locations of theelectrode, the separator, and Li+-Nafion (if Li+-Nafion was appliedto the cell of interest). The electrode location is expected to show acharacteristic Li2S signal with a high edge height intensity, whereasthe separator location is not expected to show a signal in the sulfurabsorption edge, because the starting material, Li2S, is not soluble inthe electrolyte. The lithiated membrane served as an additional pointto verify cell alignment, as the SO3

− end groups could also be detectedin the studied energy range.

A Li2S reference was measured by preparing an ex-situ sample ofone fabricated Li2S electrode facing the aluminized Kapton window.An S8 reference was measured using the synthesized S8/Vulcan car-bon composite diluted with additional Vulcan carbon to 0.05 wt-% ofS8 to minimize self-absorption. As a representative of a polysulfidereference, a mixture of Li2S and S8 powders to give a 50 mM Li2S4

solution was added to the DOL:DME based electrolyte. The measuredXAS spectra of these three references were used to qualitatively an-alyze the X-ray absorption spectra obtained during operando studies.Quantitative analysis of the spectra was not possible due to the pres-ence of self-absorption effects at sulfur concentrations above 30 mM(30 mM S = 3.75 mM S8).18,39

Results and Discussion

In a recent mechanistic study of Li-S batteries, Gorlin et al. havedemonstrated that the overall amount of sulfur containing species (S8

and polysulfides) in the separator phase is significantly lower duringthe first charge of Li2S than in the subsequent discharge and chargeprocesses.32 This observation is visualized in Fig. 1, which showsthe voltage curve in panel a) and the raw absorption at an energyrepresenting the edge height at 2487.3 eV in panel b). The signalat this energy is proportional to the total amount of S-atoms present(dissolved S8 and Sx

2−) in the electrolyte. During the first charge,the concentration of sulfur containing species noticeably increaseswithin 50% of the charge but only to a much lower level comparedto subsequent cycles. A single peak at the S8-characteristic energyof 2472 eV is seen in the X-ray absorption spectrum taken afterapproximately 40% of the 1st charge (Fig. 1c, blue line). The solepeak indicates that only S8 species and not polysulfides are presentin the electrolyte. Once the first discharge begins, the concentrationof sulfur containing species increases dramatically (Fig. 1b), and thesame high amount of sulfur containing species is observed in the

Figure 1. a) Voltage curve over time for C/10 charge (blue) of an Li2S cathodeas starting material followed by a C/5 discharge (green) and subsequent secondcharge (orange); b) raw absorption intensity at 2487.3 eV in arbitrary units overtime for the sequences described in a), showing the correlation between the lowlevel of sulfur species in the electrolyte and the overpotential of the first charge;c) two representative normalized spectra, which were taken in the separatorlocation from the first (blue) and second charge (orange) after approximately40% of the overall charge; the dashed lines show two normalized referencespectra: 1. reference solution containing a stoichiometric mixture of Li2S andS8 in DOL:DME to give Li2S4 and 2. S8 powder diluted to a concentration of0.5 wt-% (upshifted for clarity) [All data in this figure is a representation ofthe data set published by Gorlin et al.32].

electrolyte during the second charge. An additional shoulder featureis seen at an energy of 2470.5 eV in the X-ray absorption spectrumtaken after about 40% of the second charge. The appearance of thisadditional feature indicates that during the second charge, polysulfidespecies are present in the electrolyte. However, the reason for the lackof Sx

2− species during the first charge of Li2S remains unclear becauseit is well known that S8 and Li2S undergo a chemical reaction in aproticelectrolytes to form polysulfides.40 Consequently, one would expectthat in a Li-S battery assembled in a discharged state, Li2S (startingmaterial) and sulfur (reaction product observed by XAS (cf. Fig. 1))would also react chemically to produce polysulfides. In the following,we present a series of experiments that focus on the chemical reactionbetween S8 and Li2S and explain its role in the charging mechanismof Li-S batteries.

Chemical reaction between Li2S and S8.—To investigate thechemical reaction between S8 and Li2S, our first experiment wasdesigned without the added complexity of the electrochemical envi-ronment by omitting the counter electrode. For a chemical processto occur, it is necessary that both compounds are in close contactwith each other, which was achieved by adding an S8/Vulcan-carbon-containing interlayer that was coated on a Celgard-separator into thestandard Li2S-Li-cell. The two active materials are inserted in a waythat they are conductively connected and facing toward each other.The mass-ratio of Li2S present in the electrode and S8 coated onto

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Journal of The Electrochemical Society, 165 (7) A1288-A1296 (2018) A1291

Figure 2. a) Normalized X-ray absorption spectrum of the separator spotadjacent to the S8 interlayer in the Li2S-Sinterlayer-separator configuration (noanode); XAS spectra show that the chemical reaction between S8 and Li2S cantake place to form polysulfides; the dashed line shows a normalized spectrumof a reference solution for Li2S4 (upshifted by 0.2 for clarity); b) voltage curveover the specific capacity for the first charge at a C-rate of C/10 (based on massof Li2S) of the Li2S-Li cell with the S8 containing interlayer (loading Li2S:1.5 mgLi2S/cm2, S8: 2.5 mgS/cm2; electrolyte: 1 M LiClO4, 0.5 M LiNO3 inDOL:DME, 1:1 vol:vol) after an OCV period of 30 minutes (blue) and 24 hours(black). The vertical blue line marks the theoretical capacity; independentof OCV period, the cell only charges at a potential below 3.0 V for about50 mAh/g.

the interlayer is roughly 3:5. In contrast to Li2S, solid S8 can read-ily dissolve in the solution and diffuse into the Li2S/Vulcan-carbonstructure. Once the Li2S particles and soluble S8 are in direct prox-imity, a chemical formation of polysulfides is expected and with thispresence of polysulfides, an enhancement of the charging process isanticipated.

The described setup without lithium counter electrode (Li2S elec-trode, S8 coating as interlayer, glass fiber separator with electrolyte)was applied in an XAS cell (see inset of Fig. 2a), and XAS wasmeasured in the glass fiber separator location to probe the electrolytecomposition. The resulting spectrum, which was taken approximately1 hour after the cell assembly, can be seen in Fig. 2a. It has two well-defined peaks, one at 2470.5 eV and the other at 2472 eV, which are thesame two peaks that are present in the polysulfide reference (stoichio-metric mixture of Li2S and S8 to give 50 mM Li2S4 in DOL:DME)given by the dashed line. Thus, it can clearly be observed that thechemical reaction between Li2S and S8 does occur under the pre-sented condition and that polysulfides are produced in this reaction.

To see if this setup can benefit the Li2S charge, we tested thestacked electrodes versus a lithium metal anode (see inset of Fig. 2b).

As the open circuit voltage (OCV) time might be an important factorregarding dissolution, diffusion, and chemical processes, we testedthe cells at two different OCV periods of either 30 minutes or 24hours. The resulting voltage curves over the specific capacity for aC/10 charge are shown in Fig. 2b. Both the current and the mass usedto calculate the specific capacity are based on the mass of Li2S. Thevoltage curve reveals that there is only a small initial low plateau forabout 35 mAh/gLi2S after 24 hours of OCV, which confirms the resultsfrom Fig. 2a: Polysulfides must have formed chemically, and they areresponsible for a low charging voltage at the beginning of the pro-cess. Their initial presence, however, is not sufficient to maintain thelow overpotential throughout a significant portion of the charge. Anyfurther chemical reaction does not lead to a buildup of the polysul-fide concentration in the electrolyte. Furthermore, when comparingthe black (30 min OCV) and the blue curve (24 h OCV), it becomesobvious that there is not a significant difference between both OCVperiods, indicating that the time that is given to the system prior to thecharge is not the limiting factor.

At the end of the charge, a higher specific capacity is observedthan is theoretically expected when only taking the weight of Li2Sin the cathode into account. In principle, since S8 in the interlayeris in its fully charged state, it should not contribute to the specificcapacity of the cell. Nonetheless, it can dissolve in the solution andreact with the lithium metal of the anode to produce polysulfides. Theformation of these polysulfides adds active material that originatedfrom S8 interlayer and not Li2S electrode and leads to a higher specificcapacity than would be expected based purely on the added mass ofLi2S.

The experiments presented above clearly demonstrate that al-though S8 and Li2S can react chemically to produce polysulfides (cf.Fig. 2a), the accumulation of polysulfide species is suppressed in anoperating battery during the first charge. These observations have ledto the hypothesis that the anode must also be interacting with solutionspecies and hindering their concentration buildup during operation.In a typical chemical reaction between S8 and Li2S performed in avial, the resulting polysulfide intermediates simply accumulate in theaprotic electrolyte. In a Li2S-Li cell, however, they may diffuse towardand react with the lithium anode, and thus, become unavailable forfurther chemical reactions with the cathode. Hence, a diffusion bar-rier suppressing the shuttle process should help maintain an adequatepolysulfide concentration and as a consequence allow charging thecell at a low overpotential. The subsequent experiments are aimed ata detailed analysis of this hypothesis.

Influence of the anode.—To separate the cathode compartmentfrom the anode, we have chosen to use a lithiated Nafion membrane(Li+-Nafion). The negatively charged sulfonate end group (SO−

3 ) ofLi+-Nafion allows Li+-ion transport and stops the diffusion of neg-atively charged polysulfides. Other than lowering the diffusion coef-ficient of S8, it does not create an extra electrostatic barrier for theneutral dissolved S8 molecules. The membrane can be implementedin our custom-built spectro-electrochemical cell, leaving all other ex-perimental conditions unchanged.

The voltage profile of the first charge of Li2S at a rate of C/10 inthe presence of the Li+-Nafion is shown in Figure 3a. The chargingvoltage (black line) is at a low overpotential throughout the entireprocess, contrasting the results from the cell configuration without thediffusion barrier (blue line), which exhibits a charge at a high over-potential. Because Li+-Nafion stops polysulfide diffusion, we canpostulate that the containment of polysulfides at a significant con-centration in the cathode compartment is the essential aspect for theimproved charging behavior. These polysulfides can react chemicallywith both Li2S (starting material) and S8 (oxidation product) and thusfacilitate the conversion of insulating and insoluble Li2S to solutionintermediates that can be subsequently electrochemically oxidized.27

The low overpotential in the setup with a diffusion barrier is consis-tent with recently published data from Wang et al., who utilized anLATP ceramic plate to separate cathode and anode compartments andused commercial micrometer-sized Li2S particles suspended in the

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A1292 Journal of The Electrochemical Society, 165 (7) A1288-A1296 (2018)

Figure 3. a) Voltage curve over the specific capacity for the first chargeof commercial Li2S with Li+-Nafion interlayer (black) showing a chargingplateau at 2.4 V in comparison to the charging curve without the interlayer(blue) showing a charging plateau at a higher voltage of 3.25 V (loading Li2S:1.5 mgLi2S/cm2, electrolyte: 1 M LiClO4, 0.5 M LiNO3 in DOL:DME, 1:1vol:vol); b) voltage curves over the specific capacity for two repetitions of thefirst charge of Li2S with the Li+-Nafion interlayer for different C-rates: C/2(black), C/5 (orange), C/10 (green); with increasing C-rate the duration of theplateau without overpotential decreases.

electrolyte of the cathode compartment as the active material.41 Sincethe chemical conversion of Li2S to electrochemically active solutionintermediates and the electrochemical oxidation of these intermedi-ates occur simultaneously during the charging process, it is necessaryfor the chemical reactions to be faster than the electrochemical reac-tions, for the entire charge to occur at a low overpotential. Hence, byvarying the rate of the electrochemical reaction via the charging rate,it is possible to explore the kinetics of the chemical reaction. As ex-pected, the charging curves at different rates show a clear dependenceof the voltage plateaus on the applied C-rate as shown in Fig. 3b.While the charging curve at a rate of C/10 remains at a low potentialfor most of the charge, faster rates show an earlier sudden increase to ahigher potential. Specifically, the low overpotential can be maintainedfor about half the charge at a rate of C/5, but, at a rate of C/2, it isonly possible to charge at a low overpotential for a small fraction ofthe process. These observations indicate that the chemical processesin DOL:DME based electrolyte are relatively slow and become limit-ing at high current rates. Fig. 3b shows two representative repetitionsof the experiment to demonstrate the repeatability of the experiment.The slight difference in repetitions is typical to lithium-sulfur batteriesand is likely explained either by small differences in how the cells areassembled, which can lead to a different amount of electrolyte presentin electrode pores or the separator, or by differences in the length of

time between the end of cell assembly and the beginning of the elec-trochemical experiment. During the OCV time, the electrode is wet-ted by the electrolyte and any impurities present in Li2S electrode42

can react with Li2S resulting in different starting conditions of thecell.

To further investigate the relative concentration and the composi-tion (S8 and S2−

x ) of sulfur species that form during the charge of Li2Sin the presence of Li+-Nafion, we have performed spatially resolvedoperando XAS measurements at a charging rate of C/5, where roughlythe first half of the charge occurs at a low overpotential and the secondhalf of the charge occurs at a higher overpotential. These experimentalconditions have the advantage that the presence or absence of polysul-fide species can be probed during one single charging process, whichhas regions of both high and low overpotentials for a sufficient amountof time. Using our custom-built spectro-electrochemical cell, it waspossible to not only characterize species forming in the cathode elec-trode structure, but also the species that were present in the separatorof the cathode compartment.32,43

Prior to the operando experiment, the alignment of the cell com-ponents at OCV conditions was determined via X-ray fluorescencemapping of the sulfur signal. Fig. 4a shows the obtained fluorescencemap of sulfur intensity. The map is positioned next to a schematicrepresentation of cell components. The region with the highest sulfurintensity corresponds to Li2S cathode. Adjacent to the cathode, thereis a region with no sulfur content corresponding to the glass fiber sep-arator. Next to the glass fiber separator, another region with significantsulfur intensity can be detected. It corresponds to the SO3

− end groupsof the Li+-Nafion interlayer. The visualization of the cell alignmentallowed us to find appropriate spots for the XAS measurements duringthe charging process of the cell. In particular, the XAS measurementsfocused on two spots: Li2S electrode and the separator of the cathodecompartment.

The voltage curve collected during the operando experiment is pre-sented in Fig. 4b. As expected, at a C-rate of C/5 the charging potentialremained at 2.5 V for about 60% of the charge and then increased toabove 3.0 V for the remaining 40% of the charge. Throughout theexperiment, XAS spectra were continuously measured in the two lo-cations of interest in an alternating sequence. Figs. 4c and 4d presentrepresentative spectra corresponding to the initial OCV spectra (i),spectra corresponding to the charging potential of 2.5 V (ii), spectracorresponding to charging potential of greater than 3.0 V (iii), andspectra toward the end of the experiment (iv). The exact points of thecharge at which the shown spectra were recorded are indicated bytriangles in Fig. 4b, with a downward facing triangle correspondingto the spectra taken in the cathode electrode (�) and upward facingtriangle corresponding to the spectra taken in the cathode separator(�). The OCV spectrum of the electrode (Fig. 4c, spectrum (i)) showsa clear Li2S characteristic with the two peaks at 2473 eV and 2476 eV.As the charge proceeds, the electrode spectrum (Fig. 4c, spectrum(ii)) still shows the presence of Li2S, as evidenced by the presence ofthe second peak at 2476 eV, but it also displays a newly developedS8 feature, as evidenced by an increase in absorption intensity of thefirst peak and its shift from 2473 eV to 2472.5 eV. An even strongerS8 signal and a weaker Li2S signal can be detected later in the charge(Fig. 4c, spectrum (iii)), after the charging potential increases above3.0 V. Towards the end of the charge (Fig. 4c, spectrum (iv)), the firstpeak has shifted to the characteristic S8 energy of 2472 eV, indicating aclose to complete conversion of Li2S to S8. We note that in addition tothe main peak at 2472 eV, the last spectrum contains a small shoulderat 2473 eV. Although the presence of the shoulder can be interpretedas a remaining feature of Li2S, we believe that a major part of thefeature is an artifact of self-absorption. In previous XAS studies, thesame shoulder feature is present in concentrated sulfur samples thatexperience self-absorption, but is absent in dilute sulfur samples, inwhich an undistorted sulfur spectrum is measured.18,44

Fig. 4d shows equivalent data for the separator of the cathodecompartment. The OCV spectrum at point (i) is omitted because theoverall S-signal in the separator is at a very low level at the start ofthe experiment (as shown later). As the charging process starts at a

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Figure 4. a) X-ray fluorescence map showing the relative sulfur intensity inside the cell; b) voltage curve over the specific capacity of the measured cell at aC-rate of C/5 with Li+-Nafion interlayer (loading Li2S: 1.5 mgLi2S/cm2, electrolyte: 1 M LiClO4, 0.5 M LiNO3 in DOL:DME, 1:1 vol:vol). The symbols on thevoltage curve signify where (� = cathode electrode, � = cathode separator) and when the XAS spectra i)-iv) in c) and d) were taken; c), d) normalized X-rayabsorption intensity over the incident X-ray energy from the cathode electrode (c) and cathode separator (d) locations; the spectra are upshifted by 0.75 for clarity.The electrode spectra show a conversion from Li2S to S8, while the separator spectra demonstrate the presence of polysulfides during the charging process at a lowoverpotential and their absence during the charging process at a high overpotential; panel c) and d) contain reference spectra for S, Li2S4, and Li2S in the dashedlines.

low overpotential, the separator spectrum develops two clear peaksat 2470.5 eV and 2472 eV in the sulfur absorption edge (Fig. 4d,spectrum (ii)), indicating the presence of polysulfide species.16 Atboth points (iii) and (iv) of the charge, where the charging potentialexhibits a high overpotential, the characteristic polysulfide feature at2470.5 eV is not present anymore, indicating that polysulfide specieshave been depleted. The remaining characteristic S8 feature suggeststhat only sulfur remains in the separator.

To determine how the relative concentration of sulfur species inthe separator of the cathode compartment changes throughout thecharging process, XAS spectra depicting the raw X-ray absorption(without the normalization to an edge step of 1) are shown in Fig. 5for the same experiment as that presented in Fig. 4. For clarity, thespectra are presented in two different panels: The spectra recordedduring the first 60% of the charge (charging potential of 2.5 V) areshown in Fig. 5a, while the spectra from the remaining 40% of thecharge are presented in Fig. 5b. The first spectrum (a) in Fig. 5a wasmeasured at OCV conditions, and therefore, there was not a significantS-signal detected in that spectrum. Starting with spectrum (b) andcontinuing until spectrum (j), the overall sulfur signal is continuouslyincreasing, as is evident by a growing edge height (absorption intensity

evaluated at 2487.3 eV). The rise in the overpotential at 700 mAh/gLi2S

can be directly correlated with a change in X-ray absorption spectra.In comparison to spectrum (j), spectrum (k), which was measuredduring the initial rise in the charging potential to above 3.0 V, hasboth a significantly lower edge height and a significantly smallersignal at 2470.5 eV, the characteristic peak energy of polysulfides.As the charge continues above 3.0 V, little change is seen betweenspectra (l) through (p). To visualize the observed trend better, theraw absorption intensity above the edge at 2487.3 eV is plotted versuscapacity in Fig. 5c (orange symbols). From the figure, it is clear that theincrease in sulfur intensity (sum of S8 and Sx

2−) during the first 60%of the charge is linear, and that the increase in the charging potentialcoincides with a drop in the sulfur intensity (sum of S8 and Sx

2−).Furthermore, as already discussed in the context of Fig. 4c, the risein the charging potential to above 3.0 V can also be correlated to thesudden disappearance of the polysulfide feature. Thus, the obtainedoperando results in the presence of Li+-Nafion are entirely consistentwith the observations made in the conventional cell in which solutionspecies can freely diffuse between the cathode and the anode.32 In bothexperimental setups, the charge of Li2S occurs at a high overpotential,when polysulfides are not detected in the separator. When polysulfides

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A1294 Journal of The Electrochemical Society, 165 (7) A1288-A1296 (2018)

Figure 5. Raw X-ray absorption intensity over the incident X-ray energy ofthe Li+-Nafion experiment presented in Fig. 4. Panel a) focuses on pointsa-j and panel b) on points j-p, while j is repeated for comparison; panel c)depicts the same voltage curve (black line) from Fig. 4a and visualizes boththe portion of the charge at which points a-p were generated (blue circles) andthe absorption intensity at 2487.3 eV (measure of sulfur concentration) as afunction of specific capacity (orange symbols).

are detected in the separator (the first 60% of the C/5 charge in thesetup with Li+-Nafion or the second charge in the conventional setup),the charge process occurs at ≈2.5 V.

A simple additional electrochemical experiment can furtherdemonstrate the impact of polysulfide intermediates on the charg-ing process. In this experiment, the Li+-Nafion is kept in the cell, butthe standard electrolyte in the cathode compartment is exchanged witha catholyte containing 50 mM of Li2S8 (corresponding to 400 mM ofS), and the charging rate is increased from C/5 to C/2. To estimatethe expected contribution of the added polysulfides to the capacityof the first charge, we assume that the separator can only contain avolume of 28 μL (volume of separator: 200μm × 11mm × 14mm;porosity: 0.9 in uncompressed state); if normalized by Li2S loadingof 1.5 mgLi2S this corresponds to 50mAh/gLi2S, which is negligible

Figure 6. a) Voltage curve over the specific capacity for the first charge of Li2Sat a C-rate of C/2 with Li+-Nafion interlayer (loading Li2S: 1.5 mgLi2S/cm2,electrolyte: 1 M LiClO4, 0.5 M LiNO3 in DOL:DME, 1:1 vol:vol); two casesare shown, with black curve representing the standard case and blue curverepresenting the case in which 50 mM of Li2S8 was added to the cathode com-partment. While the chemical processes in the standard case become limitingcausing the appearance overpotential, a sufficient polysulfide amount addedto the cathode compartment leads to a charging process at a low overpoten-tial; b) voltage curves of the first charge of Li2S with the lithiated Nafioninterlayer for different additives (cathode compartment only) with a fixedOCV period of 40 minutes; orange curves represent 10 mM S8, blue curvesrepresent 10 mM Li2S8, and black curves represent the control case withno additive; S8 and Li2S8 have a similar enhancing effect on the chargingprocess.

compared to the overall charge capacity of 1165 mAh/gLi2S. The re-sults of the experiment are shown in Fig. 6a. As seen in the figure,in the presence of a 50 mM of Li2S8 solution, a major portion of thecharging process can occur at a potential below 3.0 V. In the absenceof 50 mM Li2S8, however, almost the entire charging process occursat a potential above 3.0 V (see Fig. 3b). Due to this improved chargingbehavior in the latter configuration, the addition of a significant poly-sulfide concentration and restriction of polysulfides to the cathodecompartment can facilitate oxidation of Li2S to S8 at a low overpo-tential, even at charging rates as high as C/2. We note that when thesame polysulfide concentration is added in the absence of Li+-Nafion,in which case the added polysulfides are not restricted to the cathodecompartment, the main portion of the charge at the charging rate ofC/2 proceeds at a potential higher than 3.0 V. This case is illustrated inFig. A1 of Appendix A.

Comparison of S8 and Li2S8 as redox mediators.—It would beinteresting to conduct a similar experiment with the addition of 50 mMS8 to evaluate the relative effectiveness of sulfur and polysulfides asredox mediators. In principle, both species can mediate the charging

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process of Li2S by bringing it into solution, as shown by Equations 1and 2.

Li2Ssolid + x − 1

8S8,solution → 2Li+solution + S2−

x,solution [1]

Li2Ssolid + y − 1

8 − yS2−

8,solution → 2Li+solution + 7

8 − yS2−

y,solution

2 < y < 8 [2]

Unfortunately, the experiment with 50 mM S8 is not possible ina standard DOL:DME electrolyte, because S8 is soluble only up toa concentration of ≈10 mM of S8 or ≈80 mM S. Nonetheless, itis experimentally feasible to add 10 mM of either S8 or Li2S8 andcompare the relative effectiveness of the two types of sulfur species.Fig. 6b shows the electrochemical response after such an additionof the two types of species (total S concentration equals to 80 mMin both cases). We note that because the added Li2S8 is expected tocontribute only 10 mAh/gLi2S to the capacity of the first charge (seecalculation in the previous section), this contribution can be neglectedin the analysis. Experimental data was collected in the cell setup withLi+-Nafion, using C/2 charge rate based on the weight of Li2S inthe cathode, and each experiment was conducted twice to generateduplicate data. The orange curves represent the cells with 10 mM S8,the blue curves represent the cells with 10 mM Li2S8, while blackcurves represent the cells without any additives. The OCV periodbetween each cell assembly and the beginning of the charging processwas controlled to be 30–45 minutes. From the figure, it is seen thatsimilar behavior was observed in the presence of either Li2S8 or S8.Specifically, in both cases the portion of the charging process thatoccurred at a lower overpotential was approximately doubled relativeto the control experiment (black lines). This result indicates that bothspecies can equally mediate the charging process, when present in thesame concentration. However, since the solubility of S8 is significantlylower than that of Li2S8, it is less effective as a mediator during theoperation of the battery.

Equations 1 and 2 show the chemical reactions that need to occurto bring Li2S particles into solution as polysulfides, thereby facil-itating the charging process of Li-S batteries.33,45 This dissolutionstep has already been proposed by Marinescu et al. to be the “bottle-neck” of the charge.31 They demonstrated via computational studiesthat the characteristic shape of the charging curve can only be sim-ulated when including a gradual dissolution of Li2S. The results ofour study provide concrete experimental evidence to their hypoth-esis, as we show that the charge of Li2S can only proceed with alow overpotential when solution species, which are facilitating thechemical reaction with Li2S, are present. Therefore, it can be statedthat solution species and, in general, the chemical reactions that aremade accessible through these species are essential to a charge at alow overpotential. In this regard we note that if the cell is broughtinto conditions where the chemical steps are suppressed by either alack of solution species, by limiting the reaction time, or by compet-ing reactions of solution species at the anode, the charging processrequires application of a higher overpotential. Alternatively, if thechemical reactions are facilitated, for example, by decreasing the par-ticle size of the active material and thereby increasing its surface area,then the charging process requires a lower overpotential. This phe-nomenon has already been demonstrated in a number of studies inliterature.33,45,46

Conclusions

In this study, we used a combination of electrochemical and spec-troscopic experiments to generate a detailed understanding of chemi-cal processes occurring during the charge of lithium-sulfur batteries.We found that the primary oxidation pathway of Li2S is through achemical reaction step involving either S8 or Sx

2− to form electro-chemically active solution species. These solution species could thenbe electrochemically oxidized or they could diffuse toward the anode,

where they would react with the lithium metal. When this parasiticreaction at the anode was eliminated through an introduction of apolysulfide diffusion layer (e.g. Li+-Nafion interlayer), Li2S could becharged at a low overpotential using a typical charging rate of C/10.As the rate was increased to C/2, however, the charge could pro-ceed at a lower overpotential only when a significant concentrationof polysulfides (≈50 mM) was added to the separator of the cathodecompartment. Finally, the relative effectiveness of S8 and polysulfidesas redox intermediators was evaluated. Although both S8 and poly-sulfides showed similar ability to mediate the charge of Li2S whenadded at the same concentration of 10 mM, 10 mM concentration wasshown to not be sufficient to allow the charging process to occur ata low overpotential at a charging rate of C/2. Because the solubilityof S8 in the DOL:DME electrolyte is limited to only 10 mM, thisfinding clarified why solid S8 that continuously forms throughout thecharge of Li2S cannot serve as a substitute redox mediator to poly-sulfides at fast charging rates. We believe that the mechanistic insightgenerated in this work will be useful to the development of predic-tive models as well as to the rational design of lithium-sulfur batterycomponents.

Acknowledgments

The described XAS measurements were carried out at the StanfordSynchrotron Radiation Lightsource, a Directorate of SLAC NationalAccelerator Laboratory and an Office of Science User Facility oper-ated for the U.S. Department of Energy Office of Science by StanfordUniversity. Use of the Stanford Synchrotron Radiation Lightsource,SLAC National Accelerator Laboratory, is supported by the U.S. De-partment of Energy, Office of Science, Office of Basic Energy Sci-ences under Contract No. DE-AC02-76SF00515. The authors alsothank the Paul Scherrer Institut, Villigen, Switzerland for provisionof synchrotron radiation beamtime at the PHOENIX beamline of theSLS, Dr. T. Huthwelker and Dr. C. Borca for their support during ex-periments at SLS, Dr. S. Webb for his support during the experimentsat SSRL, as well as M. Wetjen and Q. He for helpful discussions. Wegratefully acknowledge the funding by the German Federal Ministryfor Economic Affairs and Energy (funding number: 03ET6045D).Y. Gorlin gratefully acknowledges the support of the Alexander vonHumboldt Postdoctoral Fellowship and Carl Friedrich von SiemensFellowship Supplement. M. Tromp gratefully acknowledges NWO(NWO-VIDI 723.014.010) and the University of Amsterdam ResearchPriority Area Sustainable Chemistry.

Appendix A

Figure A1. Voltage curve over the specific capacity for the first charge of Li2Sat a C-rate of C/2 without Li+-Nafion interlayer, but with the addition of 50 mMLi2S8 (loading Li2S: 1.5 mgLi2S/cm2, electrolyte: 1 M LiClO4, 0.5 M LiNO3in DOL:DME, 1:1 vol:vol). The low overpotential can only be maintained fora short period of time, before the charge continues at a potential above 3.0 V.This result contrasts the equivalent data obtained for the configuration withLi+-Nafion interlayer in Fig. 6a.

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ORCID

Armin Siebel https://orcid.org/0000-0001-5773-3342Yelena Gorlin https://orcid.org/0000-0002-9242-8914

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